is get a little bit more practice

constructing Lewis diagrams,

and in particular, we're going to try

to construct the Lewis diagram for formaldehyde.

Formaldehyde has one carbon, two hydrogens,

and an oxygen, CH2O.

So pause this video and have a go at it.

Try to construct a valid Lewis structure,

or a Lewis diagram for formaldehyde.

All right, now let's do this together.

Now the first step, and we saw this in a previous video,

we want to think about all of the valence electrons

for this molecule.

So we want to account,

account for the valence electrons.

Now the reason why we wanna do that

is so that while we're trying to create this structure,

we are making use of all of the valence electrons.

And to figure out how many total valence electrons we have,

we can look at a periodic table of elements.

We can see that carbon,

it's in that second row, in that second period,

so its second shell is its outer shell.

And in that shell,

it has one, two, three, four valence electrons.

So, carbon has four valence electrons.

A neutral free hydrogen atom

is going to have one valence electron,

but we have two of them here,

so it's gonna be two times one.

And then, oxygen, it also is in the second period,

and in its second shell

it has one, two, three, four, five, six valence electrons.

And so the total valence electrons in this molecule

are gonna be four plus two, which is six, plus six,

which is equal to 12 valence electrons.

Now the next step is to try to draw a structure.

Try to draw,

draw single bonds, I'll say, single bonds.

And a key question is,

what do we think is going to be our central atom?

And the rule of thumb is the least electronegative atom,

that is not hydrogen,

is a good candidate for our central atom.

So we can rule out hydrogen.

So between carbon and oxygen, we know that oxygen

is one of the most electronegative atoms,

well one of the most electronegative elements

on the periodic table of elements.

It's very close to fluorine.

And so carbon is a good candidate for the central atom.

So let's put the carbon right over here,

and then let's put these other atoms around it.

We could call them terminal atoms.

So, let's put our oxygen right over there,

and then we have two hydrogens.

Hydrogen there, a hydrogen there.

And let me draw the bonds.

So that's a single bond.

That accounts for two valence electrons.

That accounts for two valence electrons.

That accounts for two valence electrons.

So I've just used two, four, six valence electrons.

So if I subtract six valence electrons,

I am now left with six valence electrons,

six valence electrons.

So the next step is

allocate the remaining valence electrons,

trying to get to the octet rule

for atoms that are not hydrogen, and then for hydrogen,

trying to get it to have two valence electrons.

So allocate, allocate the remaining,

remaining valence electrons.

All right, so let's start with this oxygen.

This oxygen already has these two electrons

that it's sharing hanging around.

So in order to get to the octet rule, it needs six more.

So let's give it six electrons.

So, one, two, three, four, five, six.

Well I've just used up the remaining six valence electrons.

So I don't really have any more to play with,

but let's see how the other atoms are feeling.

So hydrogen here, it's able to share these two electrons

that are in this covalent bond, so it's feeling good.

It can kind of pretend that it has a full outer shell,

'cause its outer shell is just that one,

that first shell, that's filled with two electrons.

Same thing for this other hydrogen.

So at least the terminal atoms,

the oxygen and the two hydrogens,

are feeling like they have a full outer shell.

But then in the fourth step,

we're going to look at our central atom.

So, let's focus on the central atom, central atom,

and do we need more bonds,

or do we need to do something interesting here?

And what we see is that carbon,

it's able to have two, four, six electrons

hanging around it, but it would love to have eight.

Carbon would love to have a full outer shell,

so how could we do that?

Well, we could add more bonds.

Where could the bonds come from?

Well it would come from some lone pair of electrons.

Well the only lone pairs of electrons

are hanging around this oxygen.

So what if we were to take,

say, this lone pair of electrons,

and then construct another covalent bond with that?

Then, our Lewis diagram will look like this.

I will actually redraw it.

So you have your carbon,

you have your three original covalent bonds,

you had a hydrogen, a hydrogen,

and then you had your oxygen, right over here,

and now we've formed a new covalent bond, just like this,

and then you have these two other lone pairs

around the oxygen.

So let me draw that.

So, two, then another two around the oxygen.

And this is looking pretty good, because the oxygen,

it still has eight electrons hanging around,

four in lone pairs,

and then four, they're in this double bond

that it is sharing.

The hydrogens still have two electrons hanging around.

They're able to share the electrons

in each of these covalent bonds.

And now the carbon is participating in,

you could think of it as four covalent bonds,

two single bonds and one double bond,

so each of those have two electrons associated with it,

so it has eight electrons hanging around.

So this is looking really good

as a legitimate Lewis structure,

or Lewis diagram for formaldehyde.